CONTACT PROCESS Sulfuric acid is one of the most important industrial chemicals Outline three uses of sulfuric acid in industry 1. The major use of sulfuric acid in Australia is in the manufacture of fertilizers such as ammonium sulfate and superphosphate. Superphosphate is produced by reacting sulfuric acid with rock phosphate. Ammonium sulfate is produced by neutralising ammonia with sulfuric acid. 2. Production of titanium (IV) oxide from titanium minerals eg ilmenite. Titanium is an important lightweight metal used to produce strong alloys and white, opaque pigments. H2SO4 is used to leach the titanium from the minerals after mining. 3. Cleaning iron – because very corrosive used to remove the oxide layer from iron or steel before they are galvanised or electroplated. Describe the processes used to extract sulfur from mineral deposits, identifying the properties of sulfur which allow its extraction and analyzing potential environmental issues that may be associated with its extraction - Most sulphur is extracted from mineral deposits using the Frasch process. Superheated stream is pumped down the outer of 3 concentric pipes into the sulphur deposit, and since sulphur has a low melting point (119) it is readily melted. At the same time, compressed air is blown down the inner pipe, and because sulphur has a relatively low density, the air is able to force the molten sulphur up the middle pipe to the surface where it resolidifies. The insolubility of sulphur in water means that it separates from any water, leaving 99.5% pure sulphur. - Sulfur is also obtained from hydrogen sulphide in natural gas and petroleum. Incomplete combustion of H2S in a furnace produces SO2 and S. 3H2S(g) + O2(g) H2S(g) + 3S(g) + SO2(g) The mixture is cooled to condense the sulphur. - Sulfur is also released as sulphur dioxide when metal sulphide ores are smelted. Eg. ZnS(s) + O2(g) Zn(s) + SO2(g) Environmental Issues: - Sulfur is easily oxidised to sulphur dioxide or reduced to hydrogen sulfide, both of which are serious air pollutants at quite low concentrations. Care is needed to ensure that there is no inadvertent oxidation or reduction of sulphur - It is very difficult to back-fill the underground caverns left by extraction of sulphur Outline the steps and conditions necessary for the industrial production of H2SO4 from its raw materials Today most H2SO4 is manufactured by the Contact Process. Step 1 – molten sulphur (or sulphide ore eg pyrite) is combusted to form SO2. S(l) + O2(g) SO2(g) or 4FeS2(s) + 11O2(g) 4Fe2O3(s) + 8SO2(g) Step 2 – So2 gas is transferred to a catalytic converter where it is oxidised to SO3. 2SO2(g) + O2(g) 2SO3(g) + heat Conditions necessary include a pressure of 1-2 atmospheres, a small excess of O2, a catalyst of vandium (V) oxide, and temps of 400-500oC. Step 3 – SO3 is dissolved in conc H2SO4 to form oleum, H2S2O7. Water is then added to the oleum to produce H2SO4. H2S2O7(l) + H2O(l) 2H2SO4 Describe the reaction conditions necessary for the production of SO2 and SO3 - The production of SO2 is carried out in a combustion furnace. The exothermic reaction occurs quickly and goes to completion. An excess of dry air is used so that the SO2 produced is already mixed with O2 for the next step. - To produce SO3 the reaction conditions necessary include a pressure of 1-2 atmosphere, as small excess of O2 to increase the yield, temperatures of 400- 500oC and a catalyst of V2O5. The catalyst is needed to increase the reaction rate at moderate temperatures without decreasing the yield. - The conditions used are a compromise between reaction rate and equilibrium yield to produce as much SO3 as possible. Apply the relationship between rates of reaction and equilibrium conditions to the production of SO2 and SO3 The reaction to produce SO2 goes to completion. Temperature A high yield of SO3 could be achieved at low temps but the rate of the reaction would be very slow. A faster reaction rate will occur at high temps but the yield would be low. The temp chosen is a compromise between reaction rate and yield. Pressure The rate of formation of SO3 is increased by increasing the total pressure so that there are more collisions between particles. Increasing the pressure will increase the yield of SO3 because, according to Le Chatelier’s principle, the system will respond by favouring the reaction which produces fewer molecules (32). However, the actual reaction takes place at a pressure of 1-2 atmosphere because even at these pressures there is a 99.5% conversion rate. Concentration The rate of reaction is increased by increasing the concentration of reactants so there are more collisions between reactants. Increasing the concentration will also shift the equilibrium to the right, increasing the yield of SO3 and so a small excess of O2 is used because it is cost effective. Change to reaction Effect on reaction rate Effect on yield conditions Increase pressure Small increase Increase Increase concentration of Small increase Increase oxygen Remove sulphur trioxide Small increase Increase from reaction zone Increase temperature Increase Decrease Use a catalyst Increase greatly No change (but produced more quickly) Describe, using examples, the reactions of sulfuric acid acting as: An oxidising agent (put conc H2SO4 above arrows) Sulfuric acid is a moderately strong oxidising agent. An oxidising agent is a substance that brings about oxidation. Concentrated sulfuric acid oxidises many metals with the formation of sulphur dioxide and water. In this case the sulfate ion acts as the oxidant. 2- 2+ Cu(s) + SO4 (aq) + 4H+ Cu (aq) + 2H2O(l) + SO2(g) Concentrated also oxidises bromine and iodide salts to form free halogens. 7H2SO4(l) + 6KI(s) 6KHSO4(s) + S(s) + 4H2O(l) + 3I2(s) A dehydrating agent Sulfuric acid has a very strong affinity for water: it absorbs water from mixtures such as moist air and forms an aqueous solution. Concentrated sulfuric acid will remove the waters of hydration in hydrated salts. Blue copper (II) sulfate crystals are turned to white anhydrous copper sulfate by the removal of water. CuSO4.5H2O(s) CuSO4(s) + 5H2O(l) Concentrated sulfuric acid can rapidly dehydrate many organic compounds such as sugars and alcohols. Alcohols are dehydrated to form alkenes (eg converts ethanol to ethene). Concentrated sulfuric acid reacted with sucrose leaves black carbon. C12H22O11 + 11H2SO4 (aq) 12C + 11H2SO4.H2O sucrose Describe and explain the exothermic nature of sulfuric acid ionization The ionisation of sulfuric acid is exothermic, releasing lots of heat. + – H2SO4(aq) H (aq) + HSO4 (aq) + heat Sulfuric acid dissociates in two steps. – Notice that sulfuric acid is a strong acid in its first dissociation, but the HSO4 ion is a weak acid and only dissociates slightly. + – H2SO4(aq) H (aq) + H SO4 (aq) K is very large – + 2– -2 HSO4 (aq) H (aq) + SO4 (aq) K= 1.2 x 10 Identify and describe safety precautions that must be taken when using and diluting concentrated sulfuric acid - Never add water to acid – reacts with water to produce strongly exothermic reaction. This will produce a dilute solution, releasing only a small amount of heat and any splashes that occur are more likely to be water or dilute acid rather than concentrated acid. - Slowly and carefully add concentrated sulfuric acid to a large volume of water. Pour acid down the side of the container, stir frequently - Gloves, lab coat and especially safety glasses must be worn - Work near supply of running water - if acid is spilled on skin, water must be run over area for many minutes (dehydrate and loose skin) Gather, process and present information from secondary sources to describe the steps and chemistry involved in the industrial production of H2SO4 and use available evidence to analyse the process to predict ways in which the output of sulfuric acid can be maximized Today most H2SO4 is manufactured by the Contact process. Step 1 – molten sulphur (or sulphide ore eg pyrite) is combusted to form SO2. S(l) + O2(g) SO2(g) or 4FeS2(s) + 11O2(g) 4Fe2O3(s) + 8SO2(g) Step 2 – So2 gas is transferred to a catalytic converter where it is oxidised to SO3. 2SO2(g) + O2(g) 2SO3(g) + heat The output of H2SO4 can be maximised by maximising the yield of SO3 produced by carefully controlling conditions. The conditions for maximum yield include a pressure of 1-2 atmosphere, as small excess of O2 to increase the yield, temperatures of 400-500oC and a catalyst of V2O5. The catalyst is needed to increase the reaction rate at moderate temperatures without decreasing the yield. The catalyst is also in pellet form which increases its surface area. The SO2 gas is also passed over several layers of catalyst, cooling the mixture slightly in between each pass, resulting in almost complete conversion of SO2 to SO3. Step 3 – SO3 is transferred to an absorption tower and is dissolved in conc H2SO4 to form oleum, H2S2O7. Water is then added to the oleum to produce H2SO4. SO3(g) + H2SO4(l) H2S2O7(l) H2S2O7(l) + H2O(l) 2H2SO4 SO3 is not directly absorbed in water as this would vaporise the water, produce a dangerous mist of H2SO4 and result in the loss of most of the SO3, thus decreasing the yield of H2SO4. Any unreacted SO2 is recycled back to the converter so that it can be passed over the catalyst again. The final amounts of SO3 formed are sent to a second absorption tower. This can increase the conversion of SO2 to SO3 to 99.5% resulting in a max output of H2SO4. Perform first-hand investigations to observe the reactions of sulfuric acid acting as: An oxidising agent Potassium bromide + H2SO4 H2SO4 used as oxidising agent – bromine ions become bromine atoms 2Br- Br2 2KBr(s) + 3H2SO4(l) 2KHSO4(s) + Br2(g) + SO2(g) + 2H2O(l) Observation: red/orange vapour/gas forms A dehydrating agent 1.